Transition Metals

Section 1 of 45Definition of transition metals

1

Transition metal = an element with an incomplete d sub-shell, or one that can form at least one stable cation with an incomplete d sub-shell

2d-block

• d-block this definition excludes scandium and zinc, which are d-block but not transition elements

3Scandium

• Scandium it forms only Sc³⁺, [Ar] 3d⁰

4Zinc

• Zinc it forms only Zn²⁺, [Ar] 3d¹⁰

5First series

• First series titanium to copper

Section 2 of 45Electronic configurations of the first series

1Order

• Order electrons fill the lowest energy sub-shells first, and 4s fills before 3d because they overlap

2Shorthand

• [Ar] represents 1s to 3p, the argon core

Element	Configuration
Ti	[Ar] 3d2 4s2
V	[Ar] 3d3 4s2
Cr	[Ar] 3d5 4s1
Mn	[Ar] 3d5 4s2
Fe	[Ar] 3d6 4s2
Co	[Ar] 3d7 4s2
Ni	[Ar] 3d8 4s2
Cu	[Ar] 3d10 4s1

1

Exceptions

• Chromium is [Ar] 3d⁵ 4s¹, not 3d⁴ 4s².
• Copper is [Ar] 3d¹⁰ 4s¹, not 3d⁹ 4s².
• One 4s electron moves to 3d to give a half-full or full d sub-shell, which is more stable.

2Cations

• Cations remove the 4s electrons first, so Mn(III) (22 electrons) is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁴

Section 3 of 45General properties of transition metals

1Four distinctive properties

• Four distinctive properties variable oxidation states, complex ion formation, coloured compounds and catalytic activity

2

Variable oxidation states

• They can form more than one positive ion, shown by Roman numerals.
• Sodium forms only Na⁺, but iron forms Fe²⁺ (Fe(II)) and Fe³⁺ (Fe(III)).

3

Coloured ions

• [Fe(H₂O)₆]²⁺ is green.
• [Fe(H₂O)₆]³⁺ is orange-brown.

4Catalytic activity

• Catalytic activity variable oxidation states let them gain and lose electrons during catalysis,
  • and substances can also be adsorbed and activated on their surfaces

Section 4 of 45Complex ions and ligands

Complex ion = a central metal atom or ion surrounded by molecules or ions called ligands, such as aluminium(III) with six water ligands

Variety

• Variety different oxidation states allow different numbers and kinds of ligands, such as [Fe(NH₃)₆]³⁺, [Fe(OH)₆]^3- and [Fe(H₂O)₆]³⁺

Ligand = a molecule or ion that forms a co-ordinate bond with a metal by donating a pair of electrons, so it is a Lewis base and a nucleophile

Co-ordination number = the number of co-ordinate bonds to the central metal atom or ion

Ligand	Formula
Water	H2O
Ammonia	NH3
Chloride	Cl-
Cyanide	CN-
Hydroxide	OH-
Ethanedioate (ox)	C2O42-
1,2-diaminoethane (en)	H2NCH2CH2NH2

Section 5 of 45Monodentate, bidentate and multidentate ligands

Monodentate

• Monodentate one dative bond per ligand
• Examples are H₂O, NH₃, Cl^- and CN^-.
• Tetrahedral [CuCl₄]^2- and square planar [Ni(CN)₄]^2- have co-ordination number 4.
• Octahedral [Fe(H₂O)₆]²⁺ and [Co(NH₃)₆]²⁺ have co-ordination number 6.

Bidentate

• Bidentate two dative bonds per ligand, from two atoms with lone pairs
• 1,2-diaminoethane (en) and the ethanedioate ion (ox).

Multidentate

• Multidentate more than two dative bonds
• EDTA^4- is hexadentate and forms six dative bonds.

Section 6 of 45Ligand substitution: principles

Ligand exchange (substitution) = one ligand in a complex is replaced by another

Stability

• Stability the new complex is more stable than the original

Extent

• Extent substitution can be complete or incomplete

Possible changes

• Possible changes the charge, the co-ordination number and the shape of the complex

Section 7 of 45Complete substitution of cobalt(II) by ammonia

1Starting complex

• [Co(H₂O)₆]²⁺ is pink

2

With ammonia solution

• OH^- from the ammonia first partly replaces water, giving a blue precipitate of Co(H₂O)₄(OH)₂.
• Excess ammonia then replaces all the water: [Co(H₂O)₆]²⁺ + 6NH₃ → [Co(NH₃)₆]²⁺ + 6H₂O.
• The solution is pale yellow or straw-coloured (yellow-brown).

3No change

• NH₃ and H₂O are similar in size and uncharged, so the charge, octahedral shape and co-ordination number stay the same

4Excess concentrated ammonia

• Excess concentrated ammonia the solution turns brown because Co(II) is readily oxidised in air to [Co(NH₃)₆]³⁺

Section 8 of 45Incomplete substitution by hydroxide

Cobalt(II)

• Cobalt(II) dropwise ammonia or hydroxide gives a blue-green precipitate: [Co(H₂O)₆]²⁺ + 2OH^- → Co(OH)₂(H₂O)₄(s) + 2H₂O

Copper(II)

• Copper(II) blue [Cu(H₂O)₆]²⁺ gives a blue precipitate: [Cu(H₂O)₆]²⁺ + 2OH^- → Cu(OH)₂(H₂O)₄(s) + 2H₂O

Change

• Change the charge goes from +2 to 0 because two neutral water ligands are replaced by two negative hydroxides

No change

• No change the octahedral shape and co-ordination number stay the same

Further hydroxide

• Further hydroxide it causes no further change

Section 9 of 45Copper(II) with ammonia

1Precipitate first

• Precipitate first dropwise ammonia gives a blue precipitate, Cu(OH)₂(H₂O)₄

2Excess ammonia

• Excess ammonia the precipitate dissolves to a deep blue solution: Cu(OH)₂(H₂O)₄(s) + 4NH₃(aq) → [Cu(NH₃)₄(H₂O)₂]²⁺(aq) + 2H₂O(l) + 2OH^-(aq)

3Partial substitution

• Partial substitution only four of six water ligands are replaced

4No change

• No change the charge, octahedral shape and co-ordination number stay the same

Section 10 of 45Substitution by chloride ions: general features

Substitution by chloride ions: general features

Size effect

• Size effect chloride is larger than water, so fewer ligands fit around the metal ion, changing the co-ordination number and shape

Charge

• Charge it changes because chloride is charged

Reagent

• Reagent concentrated hydrochloric acid

Equations

• [M(H₂O)₆]²⁺ + 4Cl^- → [MCl₄]^2- + 6H₂O, and [M(H₂O)₆]³⁺ + 4Cl^- → [MCl₄]^- + 6H₂O

Reversal

• Reversal adding water displaces the chloride ligands and restores the aqua complex

Section 11 of 45Chloride with cobalt(II), copper(II) and iron(III)

Complex change	Colour change	Charge	Shape	Co-ordination number
[Co(H2O)6]2+ to [CoCl4]2-	Pink to blue	+2 to −2	Octahedral to tetrahedral	6 to 4
[Cu(H2O)6]2+ to [CuCl4]2-	Blue to yellow	+2 to −2	Octahedral to tetrahedral	6 to 4
[Fe(H2O)6]3+ to [FeCl4]-	Yellow to orange	+3 to −1	Octahedral to tetrahedral	6 to 4

Copper mixture

• Copper mixture the reaction is reversible, so some blue [Cu(H₂O)₆]²⁺ remains and the mixture looks green

Adding water

• Adding water pink [Co(H₂O)₆]²⁺, blue [Cu(H₂O)₆]²⁺ or yellow [Fe(H₂O)₆]³⁺ returns

Section 12 of 45The haem complex

The haem complex

Structure

• Structure a central iron(II) ion with a multidentate haem ligand in a square planar arrangement, with oxygen binding as a further ligand

Oxygen transport

• O₂ is a poor ligand, bonding weakly to Fe(II), so it releases easily into cells

Carbon monoxide poisoning

• CO is a better ligand than oxygen and binds strongly and irreversibly to iron(II).
• Oxygen can no longer be carried to the cells.
• Carboxyhaemoglobin gives a darker red colour than oxyhaemoglobin, a sign of poisoning.

Section 13 of 45The chelate effect

Chelate effect = the replacement of monodentate ligands by bidentate or multidentate ligands, which is energetically favourable

Free energy

• ΔG is negative, with ΔG = ΔH − TΔS

Enthalpy

• Enthalpy changes are usually small for ligand exchange between aqueous ions

Entropy

• Entropy it increases strongly because the number of particles rises, and this is the driving force

EDTA example

• [Co(H₂O)₆]²⁺ + EDTA^4- → [CoEDTA]^2- + 6H₂O has two reactants and seven product particles

Exam skill

• Exam skill explain it in terms of the balance between entropy and enthalpy

Section 14 of 45Linear, tetrahedral and square planar complexes

Shape

• Shape it depends on the size of the ligands and the number of co-ordinate bonds

Linear

• Linear co-ordination number 2, 180°
• Cu⁺ or Ag⁺ with two ammonia ligands.
• Ag⁺ with ammonia is Tollens’ reagent.

Tetrahedral

• Tetrahedral co-ordination number 4, 109.5°
• Typical of four chloride ligands, because large chloride ions allow only four.

Square planar

• Square planar co-ordination number 4, 90°
• Typical of cyanide ligands.
• Cisplatin is an example.

Section 15 of 45Octahedral complexes and drawing shapes

Octahedral

• Octahedral co-ordination number 6, with 90° and 180° angles

Ways to make six co-ordinate bonds

• Six small monodentate ligands such as water, ammonia, hydroxide or thiocyanate.
• Three bidentate ligands such as en or ethanedioate.
• One multidentate ligand such as EDTA^4-.

Geometry	Co-ordinate bonds	Bond angle	Example ligands
Linear	2	180°	NH3
Tetrahedral	4	109.5°	Cl-
Square planar	4	90°	CN-
Octahedral	6	90° and 180°	H2O, NH3, OH-, ox, en, EDTA4-

• Stereochemical drawings
  • A solid line is a bond in the plane of the paper.
  • A dotted (hatched) line goes behind the plane.
  • A solid wedge comes out of the paper.

Section 16 of 45Geometrical (cis–trans) isomerism

Occurrence

• Occurrence it appears in square planar and octahedral complexes that have two pairs of different ligands, as a special case of E–Z isomerism

Cis

• Cis the two different ligands are adjacent

Trans

• Trans the two different ligands are opposite

Square planar example

• Square planar example cisplatin (Z) and transplatin (E)
• Cisplatin binds to DNA in cancer cells and treats cancer.
• Transplatin cannot be used in cancer treatment.

Octahedral examples

• [Cu(NH₃)₄(H₂O)₂]²⁺, where the two water ligands are adjacent in cis and opposite in trans.
• [Ni(en)₂Cl₂]²⁺, where the two chlorides are adjacent in cis and opposite in trans.

Section 17 of 45Optical isomerism

Occurrence

• Occurrence octahedral complexes with bidentate ligands can have optical isomers

Optical isomers

• Optical isomers non-superimposable mirror images with no plane of symmetry

Difference

• Difference they rotate the plane of polarised light in opposite directions

Examples

• [Ni(en)₃]²⁺ and [Ni(en)₂(H₂O)₂]²⁺

Section 18 of 45Colour in transition metal complexes

Occurrence

• Occurrence most transition element complexes are coloured, and the metal ion can often be identified by colour

Light absorbed

• Light absorbed a coloured solution absorbs some wavelengths from the visible region

Observed colour

• Observed colour it is the complementary colour, made of the wavelengths transmitted or reflected

Example

• Example copper(II) ions absorb red-orange light, so the solution appears blue-green (cyan)

Section 19 of 45Electron promotion and d-orbital splitting

Isolated ion

• Isolated ion in an ion with no ligands, all five 3d orbitals have equal energy (degenerate)

With ligands

• With ligands the 3d orbitals split into two non-degenerate sets, separated by an energy difference ΔE

Promotion

• Promotion light of the right frequency lifts an electron from a lower 3d orbital (ground state) to a higher 3d orbital (excited state)

Energy

• ΔE = hv, or ΔE = hc ÷ λ, where h is Planck’s constant (6.626 × 10^-34), v the frequency, c the speed of light and λ the wavelength

Colour

• Colour the frequencies not absorbed combine to give the complementary colour that is seen

Section 20 of 45Factors that change the colour

Ligand

• Ligand different ligands change ΔE

Ligand effect

• Ligands with greater charge density interact more strongly, causing greater splitting.
• More energy is then needed for promotion, so a different light colour is absorbed.
• [Cu(H₂O)₆]²⁺ is light blue, and [Cu(NH₃)₄(H₂O)₂]²⁺ is dark blue, with Cu(II) in both.

Co-ordination number

• Co-ordination number it changes the orientation of ligands relative to the d orbitals, and usually changes the ligand too

Oxidation state

• A higher oxidation state gives a higher effective nuclear charge and stronger interaction with ligands.
• [Mn(H₂O)₆]²⁺ absorbs green light and appears pink.
• [Fe(H₂O)₆]²⁺ absorbs red light and appears green, while [Fe(H₂O)₆]³⁺ absorbs higher-energy blue light and appears orange.

Exam skill

• Exam skill know the factors and give equations as examples of colour changes

Section 21 of 45Visible light spectroscopy

Purpose

• Purpose it uses absorption of visible light to find the concentration of coloured ion solutions

Colorimeter

• Colorimeter it passes light through the sample, using coloured filters and a detector to measure absorption

Filter choice

• Filter choice choose the filter complementary to the solution’s colour, such as a red filter for a blue solution, so absorption is greatest

Section 22 of 45Calibration curve

1

Prepare standard solutions of known concentration.

2

Measure the absorbance of each.

3

Plot absorbance against concentration to make a calibration curve.

4

Measure the absorbance of the unknown solution of the same complex ion.

5

Read across to the curve and down to the concentration.

• Beer–Lambert law
  • at low concentrations absorbance is directly proportional to concentration, so the graph is a straight line

Section 23 of 45Why oxidation states vary

Ion formation

• Ion formation transition metals lose 4s electrons first, because electron repulsion raises the 4s above 3d once the orbitals are occupied

Common state

• +2, from the loss of the 4s electrons

Element	Common oxidation states
Ti	+2, +3, +4
V	+2, +3, +4, +5
Cr	+2, +3, +6
Mn	+2, +4, +6, +7
Fe	+2, +3
Co	+2, +3
Ni	+2, +3, +4
Cu	+1, +2

• Energy
  • the first few ionisation energies are small and close together, so oxidation states can vary easily

Section 24 of 45Trends in oxidation state

Common states

• +2 and +3

Across the series

• +3 is more stable up to chromium and +2 is more stable for later elements

High oxidation states

• High oxidation states ions of +3 and above are polarising, with high charge density, giving bonds with covalent character

Maximum state

• Maximum state it equals the total number of 4s and 3d electrons and is greatest at manganese, as in the powerful oxidising agent MnO₄^-

Section 25 of 45Redox potential: effect of pH

Dependence

• Dependence the redox potential for a change to a lower oxidation state depends on pH and the ligand

Acid

• MnO₄^-(aq) + 8H⁺(aq) + 5e^- ⇌ Mn²⁺(aq) + 4H₂O(l), E = +1.51 V, which goes to completion in excess acid as Mn goes from +7 to +2

Neutral solution

• MnO₄^-(aq) + 2H₂O(l) + 3e^- ⇌ MnO₂(s) + 4OH^-(aq), E = +0.59 V, so reduction stops at +4

Conclusion

• Conclusion a change in pH changes E

Section 26 of 45Redox potential: effect of ligand

Example

• [Ni(H₂O)₆]²⁺(aq) + 2e^- ⇌ Ni(s) + 6H₂O(l) has E = −0.26 V, and [Ni(NH₃)₆]²⁺(aq) + 2e^- ⇌ Ni(s) + 6NH₃(aq) has E = −0.49 V

Reason

• Reason ammonia is a stronger ligand than water and binds Ni(II) better, so the equilibrium lies slightly more to the left, giving a more negative E

Section 27 of 45Oxidation states of vanadium

Oxidation state	Ion	Name	Colour
+2	V2+	Vanadium(II)	Violet
+3	V3+	Vanadium(III)	Green
+4	VO2+	Oxovanadium(IV)	Blue
+5	VO2+	Dioxovanadium(V)	Yellow

Vanadate

• Vanadate vanadium(V) also exists as VO₃^- in ammonium vanadate, NH₄VO₃, a reasonably strong oxidising agent

Preparation

• Preparation adding acid to solid ammonium vanadate gives a yellow solution containing VO₂⁺

Section 28 of 45Reduction of vanadium with zinc

Principle

• Principle zinc in acid reduces vanadium(V) successively, with colour changes yellow → blue → green → violet

Mnemonic

• "You Better Get Vanadium"
  • (yellow, blue, green, violet)

+5 to +4 (yellow to blue)

• Zn²⁺(aq) + 2e^- → Zn(s) is reversed, and VO₂⁺ + 2H⁺ + e^- → VO²⁺ + H₂O is doubled.
• Overall2VO₂⁺ + 4H⁺ + Zn → 2VO²⁺ + Zn²⁺ + 2H₂O.

+4 to +3 (blue to green)

• VO²⁺ + 2H⁺ + e^- → V³⁺ + H₂O.
• Overall2VO²⁺ + 4H⁺ + Zn → 2V³⁺ + Zn²⁺ + 2H₂O.

+3 to +2 (green to violet)

• V³⁺ + e^- → V²⁺.
• Overall2V³⁺ + Zn → 2V²⁺ + Zn²⁺.

Method

• Method reverse the zinc half-equation, double the vanadium half-equation to match electrons, then add

Section 29 of 45Predicting the final oxidation state with E values

• Rule
  • a reaction is feasible if E_cell is positive, so the reducing agent’s half-equation must be more negative than the vanadium half-equation

Half-equation	E / V
V3+ + e- ⇌ V2+	-0.26
Sn2+ + 2e- ⇌ Sn	-0.14
VO2+ + 2H+ + e- ⇌ V3+ + H2O	+0.34
S4O62- + 2e- ⇌ 2S2O32-	+0.47
VO2+ + 2H+ + e- ⇌ VO2+ + H2O	+1.00

Tin

• E_cell is +1.14 V for +5 to +4, +0.48 V for +4 to +3 and −0.12 V for +3 to +2, so the final state is +3

Thiosulfate

• E_cell is +0.53 V for +5 to +4, −0.13 V for +4 to +3, so the final state is +4

Section 30 of 45Tollens’ reagent

Composition

• Composition the diamminesilver(I) complex, [Ag(NH₃)₂]⁺

Stability

• Stability it is unstable, so it is prepared when needed

Tollens’ reagent

Add sodium hydroxide to silver nitrate solution.

Add concentrated ammonia solution.

A brown precipitate forms briefly and redissolves to a colourless solution.

Tollens’ reagent

Use

• Use it distinguishes aldehydes from ketones and detects reducing sugars such as glucose

Test

• Warm a few drops of the aldehyde with Tollens’ reagent in a water bath.
• Aldehydes are reducing agents and reduce Ag⁺ to silver, giving a silver mirror on the tube.
• Ketones do not react.

Equation

• [Ag(NH₃)₂]⁺ + e^- → Ag(s) + 2NH₃(aq)

Both functional groups

• Both functional groups aldehydes and ketones contain C=O, but in aldehydes it is at the end of the chain

Section 31 of 45Redox titrations: principles

Purpose

• Purpose to find the concentration of ions that change oxidation state

Specified titrations

• Specified titrations ethanedioate (oxalate) with permanganate in acid, and iron(II) with permanganate in acid

Exam scope

• Exam scope similar titrations, such as with dichromate(VI), may be set

Manganate(VII)

• Manganate(VII) it is the oxidising agent, reduced to Mn²⁺, so the mixture must be acidified with excess acid first

Section 32 of 45Choice of acid

Dilute sulfuric acid

• Dilute sulfuric acid it is used because it does not oxidise under these conditions and does not react with manganate(VII)

Unsuitable acids

• Hydrochloric acid can be oxidised to chlorine by manganate(VII).
• Nitric acid is an oxidising agent and may oxidise the analyte.
• Ethanoic acid is weak, so [H⁺] may be insufficient.
• Concentrated sulfuric acid may oxidise the analyte.

Section 33 of 45Indicator and end point

Self-indicating

• Self-indicating potassium manganate(VII) is purple and is added from the burette

During titration

• During titration it reacts with the analyte to give Mn²⁺, which is so dilute that the solution looks colourless

End point

• End point the first permanent pale pink colour in the flask, caused by a slight excess of MnO₄^- that remains on swirling

Section 34 of 45Manganate(VII) with ethanedioate

Oxidation

• C₂O₄^2-(aq) → 2CO₂(g) + 2e^-, with carbon rising from +3 to +4

Reduction

• MnO₄^-(aq) + 8H⁺(aq) + 5e^- → Mn²⁺(aq) + 4H₂O(l), with manganese falling from +7 to +2

Manganate(VII) with ethanedioate

Multiply the oxidation half-equation by 5.

Multiply the reduction half-equation by 2.

Add them and cancel the 10 electrons.

Overall

• 5C₂O₄^2-(aq) + 2MnO₄^-(aq) + 16H⁺(aq) → 2Mn²⁺(aq) + 8H₂O(l) + 10CO₂(g)

Autocatalysis

• Autocatalysis the Mn²⁺ product catalyses the reaction,
  • so the reaction speeds up as more forms,
  • because transition metal ions can change oxidation state

Section 35 of 45Manganate(VII) with iron(II)

Method

• Method a known volume of Fe²⁺ solution is titrated with a known concentration of MnO₄^-

Observation

• Observation the purple MnO₄^- loses its colour as it reacts

End point

• End point all Fe²⁺ has reacted and the first trace of permanent pink appears, from a slight excess of MnO₄^-

Oxidation

• Fe²⁺(aq) → Fe³⁺(aq) + e^-.
• Iron goes from +2 to +3.

Reduction

• MnO₄^-(aq) + 8H⁺(aq) + 5e^- → Mn²⁺(aq) + 4H₂O(l).
• Manganese goes from +7 to +2.

Combining

• Combining multiply the oxidation half-equation by 5 to match the electrons

Overall

• 5Fe²⁺(aq) + MnO₄^-(aq) + 8H⁺(aq) → Mn²⁺(aq) + 4H₂O(l) + 5Fe³⁺(aq)

Section 36 of 45Redox titration calculations

1

Write the half-equations for the oxidant and the reductant.

2

Deduce the overall equation.

3

Calculate the moles of manganate(VII) or dichromate(VI) used.

4

Use the equation to find the ratio of oxidant to reductant.

5

Calculate the moles of reductant in the sample solution.

6

Calculate the moles in the original solution.

• Iron tablet example
  • 28.50 cm³ of 0.0180 mol dm^-3 MnO₄^- is 5.13 × 10^-4 mol, the ratio 1 : 5 gives 2.565 × 10^-3 mol Fe²⁺,
  • which is 0.143 g,
  • so 14.9% of a 0.960 g tablet is iron

Section 37 of 45Transition metal catalysts: overview

Transition metal catalysts: overview

Importance

• Importance catalysts let industrial reactions be faster and at lower temperature

Reason

• Reason transition metals and their compounds form ions with more than one stable oxidation state, and have vacant d orbitals

Types

• Types heterogeneous and homogeneous (same phase)
  • (different phase from the reactants)

Section 38 of 45Heterogeneous catalysts

Mechanism

• Mechanism the reaction occurs at active sites on the catalyst’s surface

Examples

• Examples solid vanadium(V) oxide in the Contact process, and solid iron in the Haber process

Reducing cost

• Precious metal catalysts are expensive.
• Increase the surface area.
• Coat an inert support, such as a honeycomb matrix, with a thin layer of catalyst.

Catalytic converters

• Catalytic converters they reduce exhaust pollution using finely divided platinum and rhodium on a ceramic base

Section 39 of 45The Contact process catalyst

1Step 1

• S(s) + O₂(g) → SO₂(g)

2Step 2

• Step 2 the equilibrium 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) is catalysed by V₂O₅

3

Mechanism using variable oxidation state

• SO₂ + V₂O₅ → V₂O₄ + SO₃, reducing the catalyst.
• O₂ + 2V₂O₄ → 2V₂O₅, regenerating it.

4Catalyst behaviour

• Catalyst behaviour it is not used up, it speeds up the reaction and it provides an alternative pathway

Section 40 of 45Haber process and adsorption

1Reaction

• N₂(g) + 3H₂(g) ⇌ 2NH₃(g) with iron pellets to give a large surface area

2Adsorption

• Adsorption the iron surface attracts electrons in nitrogen and hydrogen, which form temporary loose attachments to the surface

3Sites

• Sites heterogeneous catalysis takes place at active sites on the catalyst surface

Section 41 of 45Limitations of heterogeneous catalysts

Poisoning

• Poisoning tiny amounts of impurities lodge on the surface and bind to active sites, preventing further catalysis

Lead

• Lead lead additives were removed from petrol because they would poison catalytic converters, and for environmental reasons

Loss of catalyst

• Loss of catalyst the thin coating on the ceramic is gradually lost, reducing efficiency

Temperature

• Temperature converters work poorly when cold, needing about 15 minutes of running, so short journeys do little to reduce pollution

Section 42 of 45Homogeneous catalysts

Homogeneous catalysts

Principle

• Principle transition metal ions readily change oxidation state, so they can act as both oxidising and reducing agents and catalyse redox reactions

Iron

• Fe²⁺ → Fe³⁺ + e^- (reducing agent), and Fe³⁺ + e^- → Fe²⁺
  • (oxidising agent)

Iodide and peroxodisulfate

• S₂O₈^2- + 2I^- → I₂ + 2SO₄^2-
• It is energetically favourable but slow, because two negative ions repel and rarely collide.
• Fe²⁺ reduces S₂O₈^2-: S₂O₈^2- + 2Fe²⁺ → 2SO₄^2- + 2Fe³⁺.
• Fe³⁺ oxidises iodide: 2I^- + 2Fe³⁺ → I₂ + 2Fe²⁺.
• Fe²⁺ is regenerated, speeding up the reaction by an alternative pathway.

Section 43 of 45Autocatalysis

Autocatalysis = a reaction that speeds up as a catalytic product forms

Graph

• Graph concentration against time becomes steeper as the rate rises, unlike normal reactions where the rate falls

Example

• MnO₄^- with ethanedioate, 5C₂O₄^2- + 2MnO₄^- + 16H⁺ → 2Mn²⁺ + 8H₂O + 10CO₂

Mechanism

• Mn²⁺ is absent at the start and forms during the reaction.
• 4Mn²⁺ + MnO₄^- + 8H⁺ → 5Mn³⁺ + 4H₂O.
• 2Mn³⁺ + C₂O₄^2- → 2CO₂ + 2Mn²⁺, regenerating the catalyst.

Monitoring

• Monitoring a colorimeter follows the fall in purple colour, and the rate of colour change increases with time

Section 44 of 45Key terms

Term	Meaning
Transition metal	An element with an incomplete d sub-shell, or that forms an ion with one.
Complex ion	A central metal ion bonded to ligands by co-ordinate bonds.
Ligand	A molecule or ion that donates a lone pair to a metal ion.
Co-ordination number	The number of co-ordinate bonds to the central metal ion.
Ligand exchange	The replacement of one ligand by another.
Bidentate ligand	A ligand that forms two co-ordinate bonds.

Term	Meaning
Chelate effect	The favourable replacement of monodentate ligands by bidentate or multidentate ligands.
Haem complex	The iron(II) complex with a multidentate ligand in haemoglobin.
Cis–trans isomerism	Isomerism where identical ligands are adjacent (cis) or opposite (trans).
Optical isomers	Non-superimposable mirror images of a complex.
Square planar	A 4-co-ordinate shape with 90° angles in one plane.
Octahedral	A 6-co-ordinate shape with 90° and 180° angles.

Term	Meaning
Complementary colour	The colour made by the wavelengths of light that are not absorbed.
Degenerate orbitals	Orbitals of equal energy.
Electron promotion	The movement of an electron to a higher energy orbital on absorbing light.
Calibration curve	A graph of absorbance against concentration for standard solutions.
Beer–Lambert law	Absorbance is proportional to concentration at low concentration.
Redox potential	The tendency of a species to be reduced, measured in volts.

Term	Meaning
Tollens’ reagent	An ammoniacal silver(I) complex used to distinguish aldehydes from ketones.
Self-indicating	A titration in which a reactant changes colour at the end point.
Autocatalysis	Catalysis of a reaction by one of its products.
End point	The point at which a permanent colour change shows the reaction is complete.
Oxovanadium(IV)	VO2+, the blue vanadium(IV) ion.

Term	Meaning
Heterogeneous catalyst	A catalyst in a different phase from the reactants.
Homogeneous catalyst	A catalyst in the same phase as the reactants.
Adsorption	The temporary attachment of molecules to a surface.
Catalyst poisoning	Blocking of active sites by impurities.
Autocatalysis	Catalysis by a product of the reaction.
Active site	A position on a catalyst surface where reaction occurs.

Section 45 of 45Exam technique

1

Giving [Cu(H₂O)₆]²⁺ a charge change on complete NH₃ substitution — NH₃ and H₂O are both neutral, so the charge stays the same.

2

Calling scandium and zinc transition metals — they do not form ions with incomplete d sub-shells.

3

Removing 3d electrons first when forming cations — remove 4s first.

4

Saying ligands accept electron pairs — they donate pairs.

1

Saying chloride complexes are octahedral — [MCl₄] complexes are tetrahedral.

2

Saying carbon monoxide binds weakly to haemoglobin — it binds strongly and irreversibly.

3

Saying the chelate effect is driven by enthalpy — it is driven by the increase in entropy.

4

Saying transplatin is used for cancer — cisplatin is.

1

Saying the observed colour is the colour absorbed — it is the complementary colour.

2

Saying the 3d orbitals are non-degenerate in an isolated ion — they are degenerate until ligands bond.

3

Choosing a filter of the same colour as the solution — choose the complementary colour.

4

Saying Cu(II) complexes have the same colour because the oxidation state is the same — the ligand changes the colour.

1

Using hydrochloric acid to acidify a manganate(VII) titration — it is oxidised by manganate(VII).

2

Giving the end point as colourless — it is the first permanent pale pink.

3

Saying vanadium(V) is blue — it is yellow, and +4 is blue.

4

Forgetting to balance electrons before combining half-equations — match the electrons first.

1

Using the same moles of Fe²⁺ as MnO₄^- — the ratio is 5 : 1.

2

Calling iron in the Haber process homogeneous — it is a solid, so heterogeneous.

3

Saying poisoned catalysts are used up — impurities block active sites.

4

Saying autocatalysis slows as products form — the rate rises.